Chemistry. for the life and medical sciences. Mitch Fry and Elizabeth Page. second edition

Transcription

1 hemistry for the life and medical sciences Mitch Fry and Elizabeth Page second edition

2 ontents Preface to the second edition Preface to the first edition about the authors ix x xi 1 elements, atoms and electrons Isotopes Electrons Summing up Test yourself 8 Taking it further: Isotopes in biology 9 Taking it further: The periodic table 13 2 Bonding, electrons and molecules What is a covalent bond? on-bonding electrons lone pairs Pi molecular orbitals oordinate bonds Electronegativity and polar covalent bonds What effect does electronegativity have on covalent bonds? Ionic bonds The concept of the chemical bond Summing up Test yourself 28 Taking it further: The peptide bond 28 3 Interactions between molecules ydrogen bonding harge charge interactions Short range charge charge interactions ydrophobic interactions Summing up Test yourself 41 Taking it further: Solubility in water 42 Taking it further: The gas laws 47 4 ounting molecules Moles Molecular mass Moles and molarity A note on units Dilutions Percent composition solutions Summing up Test yourself 60 Taking it further: onfidence with moles 61

3 vi atch Up hemistry 5 arbon the basis of biological life The electronic structure of carbon ybridisation The tetravalency of carbon Shapes of molecules arbon in chains and rings delocalisation of electrons Aromaticity Functional groups and carbon families Summing up Test yourself 80 Taking it further: arbon structures 80 Taking it further: macromolecules 83 6 The same molecule but a different shape Isomers ptical isomerism Geometric isomerism Isomers as a problem Summing up Test yourself 95 7 Water the solvent of life Bonding in the water molecule The dissociation (auto-ionisation) of water Acids and bases Using p as a measure of acidity alculating the p of water The dissociation of weak acids and weak bases in water Buffers and buffered solutions alculating the p of buffer systems using the enderson asselbalch equation Life in water Amino acids ontrolling cellular p Summing up Test yourself 111 Taking it further: Biological buffers Reacting molecules and energy Energy from molecules Getting molecules to react Energy, heat and work: some basic terms of thermodynamics Enthalpy Entropy Gibbs free energy and work Energy changes in biological reactions Summing up Test yourself 124 Taking it further: Free energy and metabolic pathways 125

4 ontents vii 9 Reacting molecules and kinetics Rate equations Reaction routes or mechanisms The rate-limiting step onsidering the activation energy Equilibrium The equilibrium position can change Free energy and equilibrium Free energy change is zero at equilibrium Transport mechanisms Summing up Test yourself 143 Taking it further: Biological transporters energy and life xidation and reduction alf-reactions Redox potential Free energy and redox potentials btaining energy for life What happens to this free energy? Summing up Test yourself 155 Taking it further: Further oxidation Reactivity of biological molecules Addition reactions Substitution reactions Elimination reactions Reaction mechanisms and reaction kinetics Free radical reactions Pi bonds and addition reactions Functional groups link molecules together Enzyme-catalysed reactions Summing up Test yourself 171 Taking it further: enzyme catalysis 172 answers to test yourself questions 179 appendices 1 Some common chemical formulae ommon anions and cations ommon functional groups otations, formulae and constants Glossary 193 Index 209

5 02 Bonding, electrons and molecules Basic concepts: Understanding the nature of covalent bonding is an essential prerequisite to predicting the behaviour of biological molecules. ere we explore the nature of covalent bonding and types of covalent bonds. Atoms react together to form molecules. Biological life forms are able to produce complex and extremely large molecules. Inherent in the stability of such molecules are the intramolecular forces between atoms which hold the molecules together, producing bonds which are strong and relatively stable. Such bonds may be referred to as covalent, dative covalent, polar covalent or ionic. 2.1 What is a covalent bond? A covalent bond is formed between two atoms, commonly by the sharing of two electrons, one electron being donated by each of the two atoms. onsider the hydrogen atom. Remember hydrogen has one proton in the nucleus and one electron in a 1s orbital. Using our diagrams of atomic orbitals, we can consider the hydrogen molecule, 2, as being formed from the overlap of two 1s orbitals, one from each hydrogen atom (Fig. 9). In this way, each hydrogen atom effectively has two electrons at any time. In other words, each hydrogen atom has a full 1s orbital. Furthermore, since both hydrogen atoms in this bond are identical, we can assume that the electrons are shared equally between the two atoms, producing a symmetrical covalent bond. This head-to-head merging of atomic orbitals forms sigma molecular orbitals. It is the formation of bonding molecular orbitals which results in covalent bonding.

6 20 atch Up hemistry hydrogen atoms () covalently bonded hydrogen molecule ( 2 ) Figure 9. Formation of a covalent bond in the hydrogen molecule Question ow can two electrons fill the 1s orbital of both atoms? Remember, orbitals represent spaces where there is a high probability of finding an electron, and electrons move close to the speed of light. So, statistically there is a high probability that at any one time both electrons will be associated with one or other hydrogen atom! Sigma molecular orbitals are just as easily formed between p orbitals, again in a head-to-head merging (Fig. 10). p orbitals forming a sigma molecular orbital Figure 10. Sigma molecular orbital formation between p orbitals

7 hapter 2 Bonding, electrons and molecules 21 Question Why do atoms form covalent bonds? taking it further: The periodic table (p. 13) In the hydrogen molecule, both atoms effectively gain a full 1s orbital, and therefore attain a stable state. For hydrogen this simply means filling its 1s orbital. In the case of nitrogen, this means filling its three 2p orbitals (see Fig. 4). Atoms whose outer (valence) electron energy level holds a complete set of electrons are substantially more stable than those which do not. By sharing electrons through covalent bonding, atoms are able to effectively fill their outer electron energy levels and so gain greater stability. For most of the lighter elements a complete outer or valence shell requires eight electrons (two electrons in the 2s orbital plus six electrons in the three 2p orbitals). 2.2 on-bonding electrons lone pairs In bonding, atoms attempt to attain a complete set of electrons in their valence shells as this arrangement generally leads to stability. In many molecules this means that atoms have electrons in their valence shells which are not involved in bonding to other atoms. These non-bonded electrons are known as lone pairs. Reminder Atoms share electrons in covalent bonds in order that each attains a full outer electron energy level, and hence greater stability onsider what happens when two fluorine atoms react together to form a molecule (Fig. 11). The electron configuration of fluorine is 1s 2 2s 2 2p 5 and so each atom of fluorine has seven outer electrons. When the half-filled p orbitals of the fluorine atoms, each containing a single electron, overlap, a sigma bond is formed. By sharing electrons in this way each fluorine atom now has a filled outer valence shell of eight electrons. owever, six of the electrons on each fluorine atom are not involved in bonding. They are paired together in non-bonding orbitals and are called lone pairs. Lone pairs of electrons have an effect upon the reactivity of a molecule and also upon its shape. In the water molecule the two lone pairs and two bonding pairs of electrons on the oxygen atom are arranged tetrahedrally. This is the arrangement in which the negative centres experience the minimum repulsion. The lone pairs exert a larger repulsive effect than the bonding pairs and so the hydrogen atoms are pushed together slightly and the bond angle becomes slightly less than the tetrahedral angle, 109. verall the molecule is therefore V-shaped (Fig. 11).

8 22 atch Up hemistry (a) uter level electrons in fluorine atoms (b) verlap of two fluorine p orbitals to give a sigma molecular orbital F F Two fluorine atoms. Unpaired electrons in p orbitals ( ) F lone pairs of electrons F p atomic orbitals overlap to form a sigma orbital, leaving three lone pairs ( ) of electrons on each fluorine atom (c) Lone pairs in the water ( 2 ) molecule In water the two lone pairs of electrons are responsible for the V-shape of the molecule Figure 11. Lone pairs of electrons 2.3 Pi molecular orbitals In addition to forming sigma molecular orbitals, electrons in p orbitals can also overlap in a side-to-side fashion, forming pi molecular orbitals (Fig. 12). 2p orbitals aligned side to side pi molecular orbital formed from overlap Figure 12. Formation of pi molecular orbitals

9 hapter 2 Bonding, electrons and molecules 23 Pi bonding between atoms occurs in addition to, rather than instead of, sigma bonding. When both a sigma molecular orbital and a pi molecular orbital are formed between two atoms, a double bond results. Whereas sigma bonding, shown by shorthand as a single line connecting two atoms, allows complete rotation about the bond, pi bonding, shown in shorthand as a double line between atoms (one sigma bond plus one pi bond) restricts rotation about that bond. Yes free rotation about bond o rotation about bond is restricted single bond double bond taking it further: The peptide bond (p. 28) Pi bonding has important consequences in determining the shapes (conformations) of biological molecules, and particularly protein molecules. Reminder ovalently bonded atoms can rotate freely about a sigma molecular orbital, but not about a pi molecular orbital 2.4 oordinate bonds A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple covalent bond, each atom supplies one electron to the bond, but that doesn t have to be the case. A coordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. In simple diagrams, a coordinate bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it. In Fig. 13, the nitrogen atom in the ammonia molecule is donating its pair of electrons to the empty 1s orbital of the positive hydrogen ion, a proton. A new dative covalent bond is formed in the ammonium ion, 4 +. nce formed, each of the - bonds is equivalent, irrespective of the source of the electrons.

10 24 atch Up hemistry + lone pair of electrons on nitrogen atom + ammonium ion forms coordinate bond with proton Figure 13. oordinate covalent bond formation in the ammonium ion 2.5 Electronegativity and polar covalent bonds taking it further: The periodic table (p. 13) In covalent bonds between like atoms, electrons are shared equally between the two atoms in the bond. owever, certain types of atoms are able to exert a greater pull on electrons than others. This ability to attract electrons within a bond is called the electronegativity of that element. The electronegativity of an atom is a property dependent on the size of the atom, and the degree of shielding of the positively charged nucleus by the negatively charged electrons. The section of the periodic table below (Fig. 14) shows the relative electronegativities of a number of atoms; the larger the value the greater their electronegativity. 2.2 Li 1.0 Be 1.5 B F 4.1 a 0.9 Mg 1.2 Al 1.5 Si 1.7 P 2.1 S 2.4 l 2.8 K 0.8 a 1.0 Ga 1.8 Ge 2.0 As 2.2 Se 2.5 Br 2.7 Figure 14. Electronegativity values of some elements of the periodic table In general, elements on the right-hand side of the periodic table are more electronegative than those on the left-hand side. In addition, as atoms get bigger (i.e. going down a group in the periodic table), electronegativity decreases [see Taking it further: The periodic table ]. In biological molecules, oxygen and nitrogen are particularly important; their electronegativity has important consequences in terms of the reactivity and associations of molecules.

11 hapter 2 Bonding, electrons and molecules 25 Reminder Electronegativity is a measure of the degree to which an atom draws electrons towards itself in a bond 2.6 What effect does electronegativity have on covalent bonds? In a sigma molecular orbital between carbon and oxygen, both atoms contribute one electron to the bond. xygen draws electrons towards itself because it is relatively more electronegative than carbon. onsequently, the two electrons in the bond are more likely to be found closer to the oxygen atom than the carbon atom (Fig. 15). This unequal electron distribution results in the formation of a polar covalent bond. partial positive charge partial negative charge Increased electron density associated with the more electronegative oxygen atom Figure 15. Unequal electron distribution in a polar covalent bond Reminder A polar covalent bond is a distorted covalent bond in which there is an unequal distribution of electrons This unequal electron distribution results in one end of the bond being slightly positive (electron deficient) and the other end being slightly negative, and thus a dipole is produced. Reminder A dipole is produced when a pair of electric charges, of equal magnitude but opposite polarity, are separated by some (usually small) distance Polar covalent bonds play an important role in biological molecules. They invariably form the basis of functional (reactive) groups on biological molecules. Such groups are responsible for the reactivity between molecules, their solubility in water and contribute to the intermolecular forces between molecules.

12 26 atch Up hemistry 2.7 Ionic bonds Ionic bonds occur when there is a complete transfer of electron(s) from one atom to another resulting in two ions, one positively charged and the other negatively charged. Electrons are not shared as in a covalent bond, but are lost or gained. For example, when a sodium atom (a) donates its one electron in its outer 3s electron energy level to a chlorine (l) atom, which needs one electron to fill its outer 3p electron shell, sodium chloride results. The bond between the two new ions (a + l ) is an ionic bond. When sodium loses an electron from the 3s atomic orbital, all available atomic orbitals in the n = 2 energy level are filled, containing a total of eight electrons (2s 2 + 2p 6 = 8), and the sodium ion, a +, has reached a more stable state. a a + + e 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 When chlorine gains an electron, all available atomic orbitals in the n = 3 energy level are filled, containing a total of eight electrons (3s 2 + 3p 6 = 8), and the chloride ion, l, has reached a more stable state. l + e l 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 As with covalent bond formation, the driving force in forming ionic bonds is to achieve a full, and therefore stable, outer electron energy level. In ionic bond formation it is the ion, rather than the atom, that reaches a more stable state and electrons are lost or gained, rather than shared, to achieve this state.

13 hapter 2 Bonding, electrons and molecules The concept of the chemical bond We have described the covalent bond, the coordinate bond, the polar covalent bond and the ionic bond. In reality, the intramolecular forces between atoms lie within a spectrum of bonding, with covalent and ionic representing extreme ends of this spectrum. The type of bonding which prevails depends upon the electronegativity difference between the atoms involved. Pure covalent bonds exist where the electronegativity difference lies between 0 and 0.4, polar covalent bonds where the electronegativity difference is between 0.4 and 2.1, and ionic bonds where the electronegativity difference is greater than 2.1 (Fig. 16). Y + X Electronegativity difference increasing δ + Y : X δ Y : X ionic (full charge on each atom) polar covalent (partial charges; asymmetrical distribution of electrons) non-polar covalent (symmetrical electron distribution) Figure 16. A spectrum of bonding between atoms The intramolecular forces between atoms provide the inherent stability for molecular construction. owever, central to biological life are the rapid and transient interactions that must occur between different molecules. Such intermolecular forces provide the basis for interaction and recognition at the molecular level of life, and are dealt with in hapter Summing up 1. A covalent bond is formed by the sharing of two electrons between two atoms, normally one electron being provided by each atom. 2. oordinate (dative covalent) bonds are formed where both electrons in the bond are provided by just one of the atoms. 3. Atoms form covalent bonds in an attempt to fill their outer atomic orbitals and so reach a more stable state. 4. Both sigma ( head-to-head ), and pi ( side-to-side ) molecular orbitals are possible through the merging of atomic orbitals. 5. Rotation about a double covalent bond is restricted (see Taking it further: The peptide bond ).

14 28 atch Up hemistry 6. Some atoms are more electronegative than others, which can result in the formation of polar covalent bonds. 7. Polar covalent bonds form the basis of functional groups on molecules. 8. Ionic bonding involves the gain or loss of electrons, the ions so formed achieving stability through attaining full outer atomic orbitals Test yourself The answers are given on p Question 2.1 Describe the relationship between atomic orbitals, molecular orbitals and covalent bonds. Question 2.2 Define the difference between a sigma molecular orbital and a pi molecular orbital. Question 2.3 What do you understand by a polar covalent bond? Question 2.4 What is unusual about a dative covalent bond? Question 2.5 What do you understand by the octet rule? Taking it further The peptide bond Proteins are polymers of amino acids. Amino acids are linked together by the formation of peptide bonds, between the carboxyl group of one amino acid and the amine group of another amino acid, in a condensation reaction (Fig. 17). A condensation reaction is a chemical reaction in which a molecule of water is lost. alpha carbon alpha carbon R R2 + 3 R1 R terminus Peptide bond -terminus Figure 17. Formation of a peptide bond between two amino acids in a condensation reaction

15 hapter 2 Bonding, electrons and molecules 29 The sequence of amino acids in a protein determines the primary structure of that protein. The flexibility and folding of the polypeptide chain is responsible for the specific three-dimensional shape of the protein, which is the main determinant of the structure activity characteristics of that protein. Proteins derive their name from the ancient Greek sea-god Proteus who could change shape; the name acknowledges the many different properties and functions of proteins. We have seen that there is free rotation about a single sigma covalent bond, but that rotation is restricted about a double (sigma + pi) covalent bond. The peptide bond is special in the sense that it is a partial double bond, and so rotation about the bond is restricted. This property of the peptide bond has a profound effect in determining the conformation of the polypeptide chain. The partial double bond character of the peptide bond The partial double bond character of the peptide bond (Fig. 18) is a consequence of the electronic configuration of the nitrogen atom, and of the pi bonding in the carbonyl group. carbonyl double bond peptide bond Figure 18. The peptide bond itrogen (), atomic number 7, has the electronic configuration 1s 2 2s 2 2p 3 Prior to covalent bonding, the atomic orbitals in the n = 2 energy level of nitrogen are hybridised. (ybridisation is covered further in Section 5.2.) ybridisation of atomic orbitals in the same energy level occurs in order to maximise the number of covalent bonds which can be formed. The more covalent bonds which can be formed, the more stable is the atom. itrogen is capable of a number of types of hybridisation; in the case of peptide formation the hybridisation is referred to as sp 2 hybridisation; the s refers to the 2s orbital, and p 2 to the fact that two of the three p orbitals are hybridised. We can see this using electrons in boxes diagrams (Fig. 19).

16 30 atch Up hemistry 2p x 2p y 2p z 2p z sp 2 hybridised orbitals 2s Figure 19. ybridisation of atomic orbitals in a nitrogen atom Electrons in the 2s orbital are raised to the slightly higher energy level of the 2p atomic orbitals; three of the resulting four orbitals are hybridised (sp 2 ), leaving one p orbital unchanged. Drawing the resulting atomic orbitals (Fig. 20) we can see that the three sp 2 hybridised orbitals are available to form sigma covalent bonds, leaving the p orbital above and below the plane of the page (the three sp 2 hybrid orbitals are arranged in a planar orientation). p z orbital (lies above and below this plane) sp 2 orbital sp 2 orbital sp 2 orbital sp 2 hybrid orbitals lie flat in this plane Figure 20. Spatial arrangement of sp 2 hybrid orbitals In Fig. 21 we have shown just the p orbitals in the peptide bond. The p orbital on the nitrogen atom is close enough to interact with the p orbitals on the carbon and oxygen atoms of the carbonyl group. The side-to-side merging of p orbitals forms pi covalent bonds.

17 hapter 2 Bonding, electrons and molecules 31 Figure 21. p orbitals of the peptide bond Electrons can move between the p orbitals (electrons are said to be delocalised) in the formation of pi covalent bonds. In fact, the peptide bond is a resonance structure, shown in Fig Figure 22. The peptide bond is a resonance structure So we see that the peptide bond has a partial double bond character; there is a variable amount (about 40%) of pi bonding between the nitrogen and carbon, sufficient to restrict bond rotation. The important properties and outcomes of the peptide bond can be summarised as follows. The rigidity of the peptide bond reduces the degrees of freedom of the polypeptide during folding. Due to the double bond character, the six atoms involved in the peptide bond group are always planar (Fig. 23).

18 32 atch Up hemistry flat (planar) region φ ψ R rotation only about the and bonds by angles φ and ψ Figure 23. Atoms around a peptide bond lie in a planar conformation Therefore, rotation about the and bonds by angles of φ and ψ respectively defines the shape of the polypeptide. If all psi and phi angles are the same the peptide assumes a repeating structure. For certain combinations of angles this can take the form of a helical structure (the alpha-helix) or a beta-sheet structure (Fig. 24). alpha-helix beta-sheet Figure 24. Alpha-helices and beta-sheets are common protein conformations Furthermore: The peptide bond is invariably found in the trans conformation, i.e. alphacarbon atoms are on opposite sides of the - peptide bond (Fig. 25). This avoids steric hindrance between groups attached to the alpha-carbon atoms (this is a form of isomerism; see Section 6.3).

19 hapter 2 Bonding, electrons and molecules 33 alpha trans alpha cis alpha alpha Figure 25. cis and trans conformations of the peptide bond The resonance donation of electrons by the nitrogen atom makes the carbonyl less electrophilic (electron-seeking). As a result the amide is comparatively unreactive. This is a good thing otherwise proteins would be too reactive to be of much use in biological systems. The special amino acid proline, a secondary amine, as opposed to other naturally occurring amino acids that are all primary amines, is unusual in that its amino group forms part of a rigid and planar ring structure (Fig. 26). Figure 26. Rotation is possible about a proline peptide bond Proline forms peptide bonds with other amino acids, but because of the ring structure, the peptide bond so formed lacks a partial double bond character. Thus, rotation about this peptide bond is possible. Therefore, when proline is incorporated into a polypeptide chain (Fig. 27), rotation about its peptide bond, plus lack of rotation about the phi --- bond (now part of the ring structure of proline), causes a kink in the polypeptide chain. The presence of proline inhibits alpha- and beta-chain conformations in proteins; proline is also found at the bends in polypeptide chains.

20 34 atch Up hemistry peptide bond with proline rotation allowed normal rotation about phi angle not possible in proline proline 3 3 free rotation about normal phi angle no rotation about normal peptide bond Figure 27. Peptide bonds in a polypeptide chain Thus we can see, at the level of the primary structure of proteins, how the special characteristics of the peptide bond are so important in constraining the three-dimensional conformations of polypeptide chains.